Will It Dissolve?

Congratulations! You’ve been selected to play “Will it Dissolve???” Do you have what it takes to leave with a prize?

It may not be as exciting as BlendTec’s Will It Blend, but answering questions about solubility can help you boost your scores on the AP Chemistry exam or your next chemistry test. And unlike crunching up an Apple Watch with high-speed blades of fruit and veggie destruction, you CAN try these experiments at home!

When faced with any problem featuring molecules in solution, think like a blender… and break the problem up into chunks. Many of the most important aspects of solution chemistry can be abbreviated with “SCOPE:” Solvent, Compound, Other ions, Precipitate, and Equilibrium.

Solvent: When something dissolves, the solvent works its way into the material so that every molecule or ion is surrounded by solvent molecules. Polar solvents like water have a side that has a slight positive charge (the hydrogens) and a side with a slight negative charge (the oxygen). They most easily dissolve polar or ionic compounds as well as molecules that have the option of hydrogen bonding. Nonpolar (uncharged) solvents like hydrocarbons are best at dissolving uncharged molecules.

Compound: Is this compound ionic or covalent? Ionic compounds (like salt, NaCl) will break up into their component ions (Na+ and Cl–) in polar solvents, while covalent molecules (like sugar) will remain intact as individual molecules.

Other solutes: Most solvents have a limited capacity for solutes. Once that capacity is reached, no more will dissolve (this is the secret behind rock candy). If another compound is added that has an ion in common with the first, they will compete for space in the solvent, and less will dissolve. Sometimes, heating the solution can allow more to dissolve depending on whether the reaction is endothermic or exothermic.

Precipitate: If a compound isn’t completely soluble, or there is too much of it, it will precipitate out of solution and form solids on the bottom of the reaction vessel. Sometimes, this is done intentionally to recover product crystals after a reaction. If you put the mixture containing the product into a solvent the product will not dissolve in, the resulting precipitate can be washed, filtered, and dried.

Equilibrium: Many compounds only partially dissolve in solvent, so their concentrations are dependent on a form of the equilibrium constant called a solubility product (Ksp). It’s calculated just like the equilibrium constant.

Now that you have a game plan for these problems, are you ready to play… Will It Dissolve??? If so, proceed to the next section! And remember, salts are said to be soluble if it at least 0.1 moles will dissolve per liter of water at room temperature. If the solution will be at a concentration of less than 0.001 M, it is said to be insoluble. Partially soluble solutions fall somewhere between those two extremes.

Round 1: Sucrose in water. Will it dissolve? And if so, how much?

Sucrose is found in table sugar. Does sugar dissolve in water? Yes! Sucrose molecules can engage in hydrogen bonding, so they will dissolve in water. They are covalent, so they don’t break up into ions—individual molecules remain intact. Sugar is the only solute in this problem, so there are no other molecules to worry about.

How much will dissolve? Here’s where things get interesting. The solubility product Ksp is only a constant at a particular temperature. A saturated solution of sugar water at room temperature will be unsaturated at higher temperatures—more can dissolve in the same amount of water. In the case of sugar water, the solubility more than doubles as water nears its boiling point.

If the maximum amount of sugar is added at a higher temperature and then the solution is slowly cooled, most of the sugar will stay dissolved for a short period of time. The resulting solution is called supersaturated. When the water cools enough, if there is some kind of impurity (a piece of string, a stick, a chip in the container), some of the sugar will precipitate out. If this happens slowly enough, the sugar will crystalize in very orderly patterns, resulting in what is often sold as “rock candy.”

For details on how to convert the number of grams of a compound that dissolve in water to the solubility product using stoichiometry, watch Example 2 (starting at minute 20) in this lecture.

Round 2: Will calcium fluoride (CaF2) dissolve in a 0.03 Molar solution of sodium fluoride (NaF2) and water? If so, how much?

This problem involves competition. CaF2 would dissolve in water according to the following reaction:

CaF2 (aq) → Ca2+ (aq) + 2F- (aq)

Because pure solids aren’t included in equilibrium constants and solubility products are calculated the same way, the solubility product will be:

Ksp = [Ca2+] [F–]2

But in this problem, the concentration of F– will be higher than normal because the sodium fluoride also releases F– ions when it dissolves. Because the solubility product is a constant, this means that the concentration of calcium ions has to decrease, which decreases the overall solubility of calcium fluoride. The solubility product of calcium fluoride is already pretty low (Ksp = 4.0×10-11). Given the added competition of sodium fluoride, it’s probably not going to dissolve at all for all practical purposes.

The answer here? A resounding “no.”

For a full explanation of this problem including all the calculations, check out Example 1 (minute 3) in this lecture.

Still playing along at home? Good! Remember, here on Will It Dissolve??? we are doing simplified versions of the problems you’ll find in your textbook or chemistry course. But it’s important to get a feeling for how these reactions work chemically and what is likely to happen. This will guide you as you crunch the numbers and help prevent mistakes. That said, are you ready for the final round? This one is a precipitation question.

When these solutions are mixed, lead, nitrate, sodium, and iodide ions will all be present in solution. What possible compounds could precipitate? To answer this, you’ll need to compare the solubility products of lead nitrate, sodium iodide, lead iodide, and sodium nitrate.

Most textbooks will have a table with some common solubility product values. These can be compared as long as each compound dissolves into the same number of ions. Unfortunately, some of these compounds break down into two ions while the others break down into three, so they’re not all comparable.

But the table will probably also list a few rules which can help here. Sodium salts are always soluble in water, as are nitrate salts, so lead nitrate, sodium iodide, and sodium nitrate will all dissolve.

Lead iodide, however, has a Ksp of 1.4 x 10-8, which is pretty small. If the product of the concentrations of lead 2+ and iodide is larger than the Ksp, then the excess ions will precipitate out of the solution as a solid. Given the concentrations of solutions, stoichiometry tells us that we have 0.16 Molar lead ion concentration and 0.48 Molar iodide.

Q (the reaction quotient) for this reaction is equal to:

Q = [Pb2+] [I–]2

Q calculates out to 3.7 x 10-5 for the concentrations of lead 2+ and iodide in this solution. 10-5 is three orders of magnitude larger than 10-8, so there are way too many ions for them all to dissolve. The excess will precipitate out as PbI2 crystals.

How much PbI2 will form? That depends on both the stoichiometry of the reaction and the equilibrium that results, so check out the very beginning of this lecture for all the calculations!

SCOPE out Your Solutions

There are a few other factors in solubility, including pH. But if you familiarize yourself with the basics of SCOPE, you’ll have a good intuitive understanding of how, when, and why things dissolve, giving you the tools you need to master the more complicated calculations. Your prize? A better understanding of the way the world works… and maybe some rock candy, if you can convince your teacher to help you run the sucrose experiment!